Unit A Chemistry - Archbishop O'Leary High Schoolyakiwchuk.weebly.com/uploads/5/9/7/3/59730135/unit_a_-_student... · Unit A ‐ Chemistry 2 Student Notes ... – The degree of hazard

Unit A – Chemistry This booklet belongs to _______________

Science 10 Unit A ‐ Chemistry 2 Student Notes

Unit A – Section 1.0 Section A 1.1 ‐ Safety in the Laboratory

Lab Safety (Personal conduct)

Most lab safety is common sense – ___________________________________ – ___________________________________ – ___________________________________ – ___________________________________

On Page 7 read safety rules; In Short form write why these rules might exist.

Ex. 1. by reading the instructions I won’t mix the wrong chemicals. Rule 2: _________________________________________________________________ Rule 3: _________________________________________________________________ Rule 4: _________________________________________________________________ Rule 5: _________________________________________________________________ Rule 6: _________________________________________________________________ Rule 7: _________________________________________________________________ Rule 8: _________________________________________________________________ Rule 9: _________________________________________________________________ Rule 10: ________________________________________________________________ Rule 11: ________________________________________________________________ Rule 12: ________________________________________________________________ Rule 13: ________________________________________________________________ Rule 14: ________________________________________________________________ Rule 15: ________________________________________________________________


Safety Hazard Symbols Each hazard symbol displays:

– The degree of hazard by the shape and color of the border

Yellow Triangle: Caution Orange Diamond: Warning Red Octagon: Danger

‐ The type of hazard is indicated by the symbol inside the border

See Fig A1.2 on page 8

Thesymbolinsidedescribesthetypeofhazard.

Example:

Askull(poison)insidea____________________requiresonly1gofpoisontokilla200lbman.

Askullina____________________________requires5gtokillthesameman

Askullina___________________________requires30g.


System

Inform

ation

Materials

Hazard

ous

Workp

lace


• MSDS sheets are provided by the _____________________________ • It identifies the chemical and physical hazards associated with each substance including:

• _______________________ • Boiling points • Toxicity • Health effects • ______________________ • Spill and clean up procedures.

Safety Apparel

You must always wear protective __________________

You must always wear shoes in the lab. The shoes must COMPLETELY cover the ______ (no sandals or open‐toed shoes allowed).

You may not wear ___________ clothing in the lab. This type of clothing can easily be dragged through a burner flame or beaker of ___________.

Long hair must be _______________ up or otherwise restrained.

Safety Environment NEVER put ___________ glass or metal in a _________ can!

Chemicals should be disposed of according to ________________ for each chemical. NEVER put chemicals in the trash cans.

No ________ or _________

Safe Work Habits • If you engage in unauthorized and/or careless work, you may be _____________ from the lab, either

for the day or permanently! • NEVER leave an experiment _______________.

Emergency Planning • Learn the locations of all safety equipment: fire extinguishers, eyewash stations, ____________, and fire

___________. • Should any type of accident or incident occur – be sure to tell your _______________!

Assignment: 1. Textbook pg 8 (Minds On)

Assignment: 1. Complete the WHMIS & HHDS worksheets on pgs 7 –

10 (Notepack)



1.1 WHMIS Symbols

1. What does WHMIS stand for? _____________________________________________________________________________________

2. What is the purpose of WHMIS? ___________________________________________________________________________________________

_____________________________________________________________________________________

3. How does the household product system developed by the Department of Consumer and corporate Affairs differ from WHMIS? ___________________________________________________________________________________________

____________________________________________________________________________________

4. What do all WHMIS symbols have in common in appearance? ____________________________________________________________________________________

5. In the science lab, bottles and cylinders containing chemical have the WHMIS labels shown in the table below. Indicate what class hazard each one is, what potential dangers are posed, and what precautions should be taken in handling.




1. When handling hazardous materials it is advisable to wear rubber gloves and safety

glasses.

2. Whenever possible, keep the materials in their labeled original containers.

3. Do not mix different types of hazardous materials.

4. It is best not to store flammable liquids in the home

5. Do not dispose of empty of full aerosol cans in ordinary garbage container

6. All hazardous materials (especially pharmaceutical drugs) should be made inaccessible

to small children by keeping them out of reach or under lock and key

7. Extremely dangerous and/or caustic substances should be store in a secure safe

location and locked up if necessary

8. Check home hazardous products for proper storage periodically. I possible keep only

what is absolutely essential.

9. Do not under any circumstance store hazardous materials with food products

10. All hazardous products listed in the above table should be disposed of at a proper waste

management center in your community. For pharmaceutical products, many

pharmacies will accept expired or no longer used medications for proper disposal

11. With very few exceptions such as drain cleaners and toilet bowl cleaner, do no flush

hazardous materials in the sink, toilet bowl or storm sewer

Cleaning

Products

Drain Cleaner

Toilet cleaner

Bleach

Cleaning

solvents

Disinfectants

Window cleaner

Paint and

Related Items

Paint (oil and

latex)

Lacquers

Varnishes

Stains

Paint remover

Stain remover

Tar

Preservatives

Lawn and

Garden

Fertilizers

Herbicides

Insecticides

Pesticides in

general

Lawnmower

gasoline

Auto/Battery

Auto batteries

Drycell batteries

Motor oil (new

and used)

Transmission fluid

Brake fluid

Antifreeze

Medicines

and Misc.

Various

medicines (over

the counter and

prescribed

Nail polish

Nail polish

remover

Glues

Photographic

chemicals

Some Common Hazardous Household Materials


Lab Equipment

• Beaker: used to heat and measure liquids NOT you • Test tube: used to heat or seal liquids using a stopper • Bunsen Burner: used to heat chemicals in test tubes or beakers using an open flame • Erlenmeyer flask: used to heat or seal liquids using a stopper • Test tube clamp: used to hold a test tube during heating or measuring • Scoopula: used to transfer chemicals from the reagent bottle to another measuring device • Buchner funnel: used with filter paper to separate the filtrate from the liquid • Graduated cylinder: used to measure the volume of a liquid. Keep the bumper ring near the top to

prevent breakage if it tips over. • Mortar and pestle: used to grind substances into a fine powder or pulp • Crucible tong: used to pick up and hold items that are too hot to handle • Watch glass: used to mix small amounts of powders or liquids • Pipette: used to add fluid to an experiment • Bulb: used to draw fluid into the pipette • Weigh boat: used to weigh powders on a scale • Safety goggles: protective eyewear used when conducting experiments • Hot plate: used to heat liquids • Burette clamp: fastened to a universal stand to hold test tubes and other apparatus • Medicine dropper: used to add small amounts of liquid to an experiment • Iron ring clamp: fastened to support various apparatus on the universal stand • Voltmeter: used to measure voltage • Ampmeter: used to measure frequency • Spot plate: Used to separate miniscule amounts of liquid • Tweezers: used to pick up miniscule items ex. Fibers • Funnel: used to filter • Test tube rack: used to hold test tubes in an upright position during an experiment • Test tube brush: used to clean test tubes

**See the Pgs 7 & 8 for images of some lab equipment

Assignment: 1. Read pages 3‐11 2. Do Question 2‐7, 10‐11 on page 11




Section A 1.2 – Properties and Classification of Matter

What is Science? Science is a way to look at the _____________.

Scientific Models

A model is a representation that serves to ___________ a scientific phenomenon

Models are supported by the ______________ method which test hypotheses by making

__________________ about the outcome of an experiment before the experiment is performed. The

results provide support or refutation of the hypothesis.

Types of Knowledge • Empirical – _______________________________________________________________ (operational) • Theoretical – ______________________________________________________________ (conceptual)

Interpretation • Indirect form of knowledge that builds on a concept to ______________________________

___________________________________.


Classifying Knowledge

Qualitative observation

• ____________________________________________________. • ____________________________________________________.

Ex. Mg is an odorless, silver‐gray solid that burns with bright white light.

Quantitative observation

• Involves a ______________________________________________________. • _____________ • _____________ • _____________

Ex. A 5 cm strip of Mg ribbon burned for 3 s leaving 2 mm of ribbon in the tongs.

What is Chemistry? • The study of ______________, their _____________, their _________________, and the

_____________ that can occur to substances.

Physical Properties Physical properties – __________________________________________________________________.

Ex. color, hardness, texture, phases, density, ability to conduct heat, electricity.

Physical Change – ____________________________________________________________________. Ex. Ice Melting

Physical Properties Description

Boiling Point

Melting point

Malleability

Ductility

Color

State

Solubility

Crystal formation

Conductivity

Magnetism


Chemical Properties

Chemical Properties ‐ ______________________________________________________. Ex. ability to burn, flash point, reactions with water, air, acids, heat and litmus paper

Chemical Change – ________________________________________________________.

**Note: in both Physical and Chemical changes there is __________________________

Matter

Matter is ____________________________________________________________

STUDYING THE CHANGES TO MATTER IS THE ESSENSE OF CHEMISTRY! A chemist is most interested in the __________________ and __________________ of matter and the

changes that occur in matter. EXAMPLES: 1. Digestion – biochemistry 2. Oil or Gas plastics – petrochemical engineering 3. Corrosion – chemical engineer 4. Pollution ‐ environmental biologist

Chemical Property Description

Ability to burn

Flash point

Behavior in air

Reaction in water

Reaction in acids

Reaction to heat

Reaction to blue or red litmus


Pure substances

All _____________________________________________________________ (all particles have the same chemical and Physical properties)

2 types of Pure substance Element – ______________________________________________________. (Ex. Sulfur) Compound – ____________________________________________________. (Ex. CO2, MgCl2)

**Note the 2 different types of chemical bonding

Mixtures

Combination of _______________________________. (not chemically combined) 4 types

1. Solutions – ________________________ mixture (same parts) Looks similar throughout

Ex. Soft drinks (sugar is dissolved) 2. Mechanical Mixture – ___________________________

You can see all parts clearly Ex. Soil

3. Suspension – ____________________________ The parts are in different states

Ex. Mud – soil (solid) & water (liquid) 4. Colloid ‐ ___________________________

Similar to suspension but ____ easy to separate Colloids look a lot like solutions.

If you filter both (solution and colloid) a solution will pass through but a colloid will leave behind chunks if one of the phases was a solid.

Ex. Milk, paint, ink…


Methods of Separation 1. Filtration ___________________________________________________ (mechanical mixture). 2. Distillation _________________________________________________ (homogeneous mixture). 3. Electrolysis _________________________________________

Chemical Reactions Characteristics of a chemical reaction include:

1. The production of a new substance with UNIQUE ________________________________________. 2. A gain or release of ______________.

i. Heat (thermal energy moving) exits the system – ______________________ ii. Heat enters the system ‐ _____________________

And sometimes: A phase change such as:

______________ (gas forming) ______________ (solid forming) ______________ (solid forming just not falling)

Assignment: 1. Do pages 19‐21(Notepack) 2. Do Questions 1, 4‐8 & 11 on page 17 (Textbook)


Physical Science Assignment

1. What type of property (physical or chemical is illustrated in each of the following situations?

_____ a. Water pours from a pitcher

_____ b. A lightbulb glows when the switch is turned on

_____ c. Grass grows in the spring.

_____ d. Lead belts cause divers to sink in water

_____ e. Drain cleaner unclogs sink.

2. When a white solid is heated, a silvery liquid and a greenish gas results. Is the original solid an element or a

compound? Explain.

___________________________________________________________________________________________

___________________________________________________________________________________________

___________________________________________________________________________________________

_______________________________________________________________

3. Give an example for each of the following homogenous mixtures. a. Solid/liquid

b. Liquid/liquid

c. Liquid/gas

d. Solid/solid

Read Pure Substances and Mixtures on pages 14‐15. Use this information to answer question 5‐7.

4. Compare the following chart

Matter

heterogeneous element


5. Write the definition and give two examples for each of the following.

a. Element:

_____________________________________________________________________________________

_______________________________________________________________________

b. Compound:

_____________________________________________________________________________________

_______________________________________________________________________

c. Homogenous mixture:

_____________________________________________________________________________________

_______________________________________________________________________

d. Heterogeneous mixture:

_____________________________________________________________________________________

_______________________________________________________________________

6. A. Identify from the examples given below which mixtures are mechanical and which are solutions.

State of Matter Examples

Solid/solid Soil, cement, cookies

Solid/liquid Mud puddle, orange juice

Liquid/liquid Vinegar, cooking oil, gasoline

Gas/liquid/solid Carbonated beverages, tap water

Gas/gas Air, natural gas as it comes from the ground

Mechanical Solution

B. Most solutions you come in contact with are a combination of water as a solvent and some other substances as the

solute. Can you identify a solution that is only made up of solids?

a. What is that solution called?

b. Give an example: ________________________________________________________________

_____________________________________________________________________________

7. Fill in the type of property (physical or chemical) that is described by the following information.

Property Description Examples

Can be determined when the composition of a substance changes

Burning of carbon, rusting of nails

Can be observed without changing the composition or makeup of the substance

Density of lead, boiling point of ammonia


8. Choose two examples for each of the following categories of matter using the list of examples provided on the

right.

Categories a. Heterogeneous mixtures

________________, ________________ b. Homogenous mixtures

________________, ________________ c. Elements

_______________, _________________ d. Compounds

_______________, _________________

9. Fill in the blanks to complete each of the following statements

a. When mixtures are separated into their components, the components are usually

____________________ or __________________ or both.

b. _______________________ is the separation method that could be used to obtain clear water from

muddy water

Examples Pizza Steel Magnesium Salt Sulphur Apple juice Concrete sugar


Section A 1.3 – Developing Ideas about Matter

Food Chemistry • Various ____________________________________________ have been used to enhance and preserve

our food. • The most common way food goes bad (spoils) is when bad bacteria occur in too large a number.

Methods for Preserving Food – _______________________________________

• Heating (cooking)‐ temporarily sterilizes food (kills the bacteria) • Freezing (storage) – halts the growth of bacteria

– ____________________ (beef jerky) • Salt removes water, by osmosis, from the food and the bacteria in it. This changes the

taste of the food but kills the bacteria. – ____________________ – This process uses beneficial bacteria (lactobacilli) to create a natural

preservative (lactic acid). • Sauerkraut was made because it lasted long and had vitamin C to fight off Scurvy. • ________________________ are also made this way.

– ___________________ – smoking your bacon or fish introduces antioxidants that help control the number of bacteria

Metallurgy • Copper age‐ made tools out of copper

– ____________________ (heating before being hammered) made Cu easy to manipulate and strengthened it. Stronger tools and Weapons.

– ____________________ – separating a metal from other compounds using heating. Most Copper is mixed with other compounds and was until this time unusable.

• Bronze age – Bronze is made from ______________________. It is called an ___________ (mixture of metals).

• Iron Age – started smelting Iron in 1200 B.C. – Later figured out how to add carbon to make steel.

Note: Copper to Bronze to Iron was a _________________ to making stronger metals that enabled better tools (farming) and weapons.

• People could change matter but _________________________.

Early Descriptions of Matter • Aristotle

– All matter was comprised of _________, ___________, ___________ & __________.. – There was no smallest part of matter. In other words you could cut an element in half forever

and still have a piece of that element. – Democritus challenged Aristotle on this last point saying that there were particles called

_________ that were indivisible. (Aristotle was in the in‐crowd so no one listened to Democritus and set back science for 2000 years give or take.)

• Alchemy – branch of early science that wanted to _____________________________________. (ideally lead into gold. WE can do this now actually just not cost effective)

– Although not successful they improved lab equipment and produced many useful compounds.


• Early scientists

– Robert Boyle – _____________________________ – Antoine Lavoisier – __________________________________________________________.

(Mass of reactants = Mass of Products)

• John Dalton – 5 points of Daltons theory:

1. Atoms are tiny, ________________________________________. 2. All elements are composed of _______________. 3. Atoms of the same ___________ are ________________. 4. Atoms combine in _______________ to form _____________________. 5. Chemical reactions occur by a ___________________ of atoms.

**Dalton used a solid billiard ball model.

• J.J. Thomson (electric JJ) – Hooked up materials in a ___________________________ and found electrons are ejected. – Plum pudding model of the atom electrons evenly scattered throughout the atom – Each atom consists of one large positive charge and many small negative charges embedded in

it. Net charge = 0

**Thomson used a raisin bun model.

• Ernest Rutherford (library named after him at the U of A) – Did research in Canada – Found ________________________________________________. – The center is made up of ________________ (known as the nucleus) – ___________________________________. (In orbit)

**Rutherford used a bee hive model

• Neils Bohr – Found that ______________________________________________________________ using

spectral analysis. – Bohr model of the atom shows orbits representing ____________________________________. – energy exists in small units called quanta. – electrons circle the nucleus in orbit at a fixed distance having a fixed amount of energy. – Bohr concluded that:

• ________________________________________________________________ (2,8,8,18) • Since electrons can’t fall below that lowest energy level, negatively charged electrons

________________ with the positively charged nucleus.

Line Spectra • When atoms are heated, bright lines appear called ______________________.

• An electron absorbs energy to “jump” to a _____________________________. • When an electron falls to a lower energy level, energy is emitted Electrons are ___________________________ occupying the whole space at once at different _____________________________________.


Unit A ‐ Section 2.0

Assignment: 1. Do p. 24 of the Notepack using your notes and

textbook 2. Do #1‐8 & 11 on page 25 (Textbook) 3. Do # 2, 3, 5‐12, 15‐17 on page 27 (Textbook) 4. Prepare for Section 1.0 Exam


2.1 Periodic table and Atomic Structure

PERIODIC TABLE OF ELEMENTS: ORGANIZATION Elements are arranged _______________________________

ATOMIC NUMBER = __________________________________! EXAMPLE: C = ___ protons U = ___ protons

Elements are arranged based on similar properties o LEFT and RIGHT of the “________________________”.

LEFT of the line have __________________________________. RIGHT of the line have ____________________________. (Normally opposite of metals!) Along the staircase line are ________________________________

Metals Non Metals Metalloids

______________________ ______________________ ______________________ ______________________ ______________________ ______________________ ______________________

______________________ ______________________ ______________________ ______________________ ______________________ ______________________ ______________________

______________________ ______________________

______ ______ ______ ______ ______ ______ ______

Periodic Table Arrangement Organized by _________________________ (# of protons) and _________________________

_______________ – horizontal rows

_____________________ – the vertical rows o These have similar physical and chemical properties. You must know the following 4:

_________________________ (group 1) • soft, shiny & reactive (vigorous with water). • Have one extra electron (valence electron)

___________________________ (group 2) • Shiny and reactive but not soft • Have 2 extra electrons

___________________________ (group 17) second last • Reactive non‐ metals • Missing one electron

__________________________ (group 18) last • Non reactive • Full outer shell or energy level •

Use the following website to help you study the different parts of the periodic table. http://www.dayah.com/periodic/


Atomic Theory

____________________________________________ (still retains all the properties of an atom)

Majority of the mass is located in the nucleus (nucleons = protons and neutrons)

Atomic number is the number of ______________ (for elements it is also equal to the number of ______________.) Example: Sodium = 11 Chlorine = 17 Lithium = 3

________________________________________________________________________________________________________________________________________________________________________

Number of protons and distribution of electrons give atoms their distinctive chemical and physical properties

N

14.01 7

ATOMIC NUMBER (Number of Protons)

Atomic Mass

Chemical symbol (nitrogen)

Groups (1‐18)

Periods (1‐7)

Assignment: 1. Complete pg 28 of Student Notepack


Particle Symbol Charge Mass Location

Proton (atomic number) p+ 1+ 1.7x10‐24g

Neutron n0 0 1.7x10‐24g

Electron e‐ 1‐ 9.1x10‐28g

Example: Lithium atom: Atomic number is ‘3’

Energy Levels

Electrons can have ___________________________________________

An energy level is a region of space close to the nucleus that may contain electrons.

The electrons located ________________________ have the least amount of ___________ but require _______________ to pry them lose.

The number of electrons in each energy level follows a pattern: 2,8,8 (18) The Bohr model of the atom shows the _______________________________(the close ones have less energy)


Periodic Table Assignment

Use your periodic table of elements to complete the following table

Element Name IUPAC Symbol

Atomic Number

Group Number

Period Number

Metal (m) or Nonmetal (nm),

metalloid (t)

SATP State

Family/Series Name

1. chlorine

2. magnesium

3. 30

4. N

5. 17 5

6. 79

7. 3 Alkali metals

8. thorium ‐‐‐‐‐‐‐‐‐‐

9. 12 Liquid

10. Br

11. argon

12. 11 5

13. 19

14. calcium

15. 1 Gas

16. 58 ‐‐‐‐‐‐‐‐‐‐


All the elements in a group have the same electron configuration in their outermost shells. Electrons in the outer shell that is not full are called valence electrons.

Noble gases have no _______________________________their outer shell is full. That is why they are so stable and non‐reactive.

Assignment: 1. Complete pg 30 in Notepack


Unit A—Energy and Matter in Chemical Change Unit Support

L ine Master 5

Atoms, Protons, and Electrons

Symbol Number of Protons Number of Electrons

Electric Charge

Li 3 3 0

C

F

Mg

19 0

15 0

16 0

10 0

18 0

8 0

17 17

11 11

20 20

Fe

Ni

Au

80 0

82 0

30 0


Formation of Ions Recall that an atom of any element is ______________, so the number of protons _______________

________________________________.

An ion is a an atom (or a group of atoms) that has a _____________ or _____________ electric charge.

The formation of an ion is called ionization, and is the result of an atom either gaining or losing electrons.

Cations

Cations are ___________________________________________.

They are formed when a metal atom ________________________________________ (electrons in the outermost energy level). As a result, the ion has more protons than electrons

CATS have PAWS…

Anions Anions are __________________________________________.

They are formed when a non‐metal atom accepts electrons into its outer energy level. As a result, the _____________________________________________________________________

Formation of Cations and Anions Creation of a sodium cation:

Na Na+ and 1 electron

The number of protons only changes in

nuclear reactions, never in the

sodium atom

(11 p+ and 11 e‐)

Charged sodium ion

(11 p+ and 10 e‐)


Creation of a chloride anion Cl and 1 electron Cl‐

The Octet Rule

The octet rule says that ________________________________________________________________ ____________________________________________________________________________________

Atoms gain or lose electrons so that they have the same number of electrons as the nearest noble gas.

Some metal atoms, depending on the nature of the chemical reaction, can form __________________ __________________________________________________________________________________ Example: Copper atoms will lose either one or two electrons.

These elements are called _________________ and the first charge given on the periodic table is the most common.

Neutral chlorine atom

(17 p+ and 17 e‐)

Charged chloride ion

(17 p+ and 18 e‐)

Assignment: 1. Complete pg 33‐34 in Notepack

+Cu 2+Cu


Atoms & Ions I

Symbol NumberofProtons

NumberofElectrons

ElectricChange

Li+ 3 2 1+C 6 6 0F‐ 9 10 1‐

Mg2+ K+ Li S2‐ He 18 0 8 2‐ 17 1‐ 11 0

Ca2+ 18Fe3+ 23Fe2+ 24 78 1+ 78 2+ 82 78 30 30


Atoms & Ions II

Name Symbol Mass

Number Atomic Number

Protons Neutrons Electrons Electric Charge

fluorine atoms

F 19 9 9 10 9 0

nitride ion

N3‐ 15 7 7 8 10 3‐

boron atom

14

aluminum ion

14 3+

gold ion 116 1+

40 19 0

79 35 35

19 18 1‐

16 16 2‐

Ag 110

cesium ion

77 1+

I‐ 125


The Atomic Mass Unit

Since the mass of individual atoms is so small, it is not convenient to use a unit like grams (g).

Instead, we use the ______________________ (amu).

______________________________________

Then the _________________________ of all other atoms is determined by comparing each to the mass of an atom of carbon‐12.

For example, an atom twice as heavy has a mass of 24.00 amu, and an atom half as heavy has a mass of 6.00 amu. Example:

• NITROGEN has: o __________________ o __________________

• therefore its atomic mass is equal to: _________________

• … but the atomic mass on the ______________________ is • 14.01 amu. • WHY?

o Isotopes: • elements that have the __________________________________ but a different mass

number!

Isotopes

Isotopes are different types of atoms of the same element. How are they different? o ____________________________________________________________ o ____________________________________________________________ o Only thing left is the # of ________________. So the only difference between isotopes of the

same element is how heavy they are. Isotope Notation:

Element

Symbol

Atomic

Number

Mass

Number AZ X


Example:

Three naturally occurring isotopes of carbon:

The __________________ of an element that is listed on the periodic table is called the atomic molar

mass.

This value is calculated as the ____________________ of all of the isotopes of an element, taking the

percent abundance into account.

Assignment: 1. Complete pg 14‐15 in Notepack 2. Read pages 34 – 38 in Text (except p. 37) 3. A2.1 Check and Reflect, page 39 #’s 1 – 10

The bottom number is sometimes not written because you can determine the

atomic number from the symbol.

126C 13

6 C 146C


Draw energy level diagrams for the following isotope symbols: 1. 2. 3. Write the Symbol for the energy diagrams given below: 4. 5.

12 n0

14 p+

2 e-

8 e-

4 e-

7n0

5p+

2 e-

3 e-

For neutral atoms, the number of

protons equals the number of electrons.

Excellent!!

2412 Mg

209 F

3215 P


Atomic Number and Mass Number

Name Mass Number Atomic Number Number of

Protons Number of Neutrons

carbon-14 14 6 6 8

hydrogen-1 1 1

hydrogen-2 2 1

carbon-12

oxygen-18

4 2

20 10

64 29

35 45

7 7

16 16

20 21

9 10

26 30

53 74

calcium-40

119 50

26 33

208 82

silver-108

mercury-201


Section A 2.2 – Naming Molecular and Ionic Compounds

What are Compounds? Compounds ___________________________________________________________________.

… _______________________________________________________________________“charged”…

EXPLANATION:

If an atom ______________________________________ positive. The atom _______________________________________________ negative.

The atoms are now ions and are attracted to each other.

“______________________________________________________!”

The ions “stick together” and an _____________________ is now formed.

Properties of Ionic Compounds

All ________ at room temp NaCl(s) (high melting and boiling points)

Retain _____________ shape

_____________ in water (majority) NaCl(aq)

Always _____________________ in solution

What are Ionic Bonds? An ionic bond forms when one of the 2 elements _______________ and the other ________________

_______________.

Normally when a ____________ (left of step ladder) meets with a _____________ (right of step ladder)

Once the element has _________________________ to become more stable (full outer shell) it is called an _______.

Ionic bonds form ________________________ due to the alternating +/‐ ion arrangement.

• This is a regular three‐dimensional pattern of alternating positive and negative ions producing an electrically neutral compound since electrons can not move

This is very stable arrangement, so all ionic compounds are solid at room temperature.

_________________________________________________ Strongest bond unless _____________________________________

Naming Ionic Compounds

Naming chemicals properly is a very important aspect of Chemistry. If not done properly no one will

know which substance you are talking about.

In _______ the International Union of Pure and Applied Chemistry (____________) came up with a

naming system.

**Remember Ionic compounds

take electrons.

Bonds resulting from the force of

attraction between atoms

For sodium chloride,

there is one sodium ion

for every chloride ion

(they are in a 1:1 ratio).

**Note how both the

sodium ion and chloride

ion have full outer shells.


Rules for Type 1 1) ___________________________________________________________________

2) ___________________________________________________________________. EX. Na is

sodium

3) ___________________________________________________________________

It ends in an ide ending

Ex. Chlorine goes to Chloride

Oxygen goes to Oxide

Selenium to Selenide

Phosphorus to Phosphide

• Some periodic tables have the ion names

Metals‐ use the full name and then add “ion” like this:

i) calcium ii) sodium iii) cesium iv) Ag+ v) Au vi) cobalt vii) Ge+4

Non Metals – change end to “ide”

i) chlorine ii) fluorine iii) oxygen iv) sulfur v) iodine vi) bromine vii) selenium

Writing Ionic Compounds

When writing formulas of _______________ ionic compounds the _________________ for the

elements are written in the same order as they appear in the name.

Subscript numbers are used to indicate the _____________________________ in the compound.

The charges on the ions must balance in the chemical formula, since ionic compounds are electrically

neutral.

1. Identify the ______ and their ____________.

2. Determine the ________ of charges needed to balance.

3. The charge on the metal ion _________ to become the _____________ on the non‐metal ion.

4. The charge on the non‐metal ion crosses to become the subscript on the metal ion.

5. ______________________ of subscripts in the formula.

Practice Problems:

Compound Name Ions Compound Formula

aluminum fluoride Al3+ F‐ AlF3(s)

silver sulfide

potassium iodide

zinc nitride

calcium oxide

Assignment:

1. Do the Practice Problems

on pg 43 of Text

2. Do pg 41‐42 in NotePack


ChemicalFormula

DescriptionorUse NameofCompound

e.g. Whitesolid;wettingagent calciumchloride

1. Dietarysupplementforiodine potassiumiodide

2.MgO Whitepowder

3. Antiperspirant aluminumchloride

4.NaBr InEpsonSalts

5. Whiting,aluminumore

6. Black;lithiumreactswithair lithiumnitride

7.CaO Whitepowder,quicklime

8. Whitesolidlike bariumchloride

9. Whitesolid;tablesalt sodiumchloride

10.ZnO Protectiveoxideonzincmetal

11. Photographicemulsion silverbromide

12. Magnesiumreactswithhydrogen magnesiumhydride

13. 11%ofmineralsinseawater magnesiumchloride

14. Insolderingpaste zincchloride

15. Argentite(silverore)

16. Potash(fertilizer) potassiumchloride

17. Fluorite(prettymauvecrystals)

18. Fortoningpicturesbrown sodiumsulfide

19. Preparationofhydrogen

20. Zincblend(zincore) zincsulfide


IonicCompounds‐‐‐UnivalentMetalIons

1. Ifthefollowingpairsofelementsweremixedandheated,theywouldcombineintosolidioniccompounds.Writethenameandformulaofeachcompoundformed.

Name Formula

a) silverandiodine silveriodide

b) magnesiumandoxygen

c) magnesiumandbromine

d) calciumandnitrogen

e) zincandselenium

f) sodiumandsulfur

g) bariumandphosphorus

h) aluminumandfluorine

i) potassiumandchlorine

j) silverandoxygen


2. Writethecorrectnamesforeachofthefollowingcompounds.a) _____________________________________________________

b) _____________________________________________________

c) _____________________________________________________

d) _____________________________________________________

e) _____________________________________________________

f) _____________________________________________________

g) NaI_____________________________________________________

h) _____________________________________________________

i) _____________________________________________________

j) _____________________________________________________


Rules for Type 2 (Binary Compounds)

___________________________________________________________________

Ex. Ni2+ or Ni3+, Au3+ or Au+

Note the one on the periodic table that is found on the top is the one most commonly found in

nature. Normally you will be given enough info to figure out the charge used but if not

__________________________________ but be careful.

How to figure out which Charge was

used

– In the name it can be shown in

2 ways: systematic (uses roman

numerals) or the old method

which uses ic (bigger choice) or

ous (smaller choice) attached

to the latin name of the

element.

Ionic compounds with multivalent elements must have Roman numerals after the name of the positive (metal)

ion to indicate the charge on that ion.

Compound Name Formula

Iron (III) chloride

Lead (IV) oxide

Nickel (III) sulfide

CuF2(s)

Cr2S3(s)

**Use roman numerals ONLY when the metal element is multivalent. “Choice of charge”

• Naming is the same except for the cation being either the systematic or Old school.

• Ex. Pb2+ + Cl‐ becomes PbCl2 named Lead (II) chloride

Try the following Examples:

Cu2+ + O2‐ 1. Cr3+ + S2‐

2. Mn4+ + O2‐

Ion Systematic Alternative

Iron(III)

Iron(II)

Copper(II)

Copper(I)

Cobalt(III)

Cobalt(II)

Assignment: 1. Do the Practice Problems on pg 44 of text 2. Do pg 45‐46 of NotePack


IonicCompounds‐‐‐MultivalentMetalIons

1. Ifthefollowingpairsofelementsweremixedandheated,theywouldcombineintosolidioniccompounds.Inthisworksheet,usethemostcommonionicformofthemultivalentmetalion.Themostcommonformislistedfirstintheperiodictable.Forexample,ironexistsasboth2+and3+ions,withiron(III)beingthemostcommon.

Ions Name Formula

a) ironandsulfur iron(III)sulfide

b) copperandoxygen

c) manganeseandfluorine

d) goldandnitrogen

e) chromiumandchlorine

f) platinumandphosphorus

g) nickelandoxygen

h) cobaltandbromine

i) tungstenandiodine

j) manganeseandsulfur


2. Writethecorrectnameforeachofthefollowingcompounds.Thechargeonthemultivalentionisnotgivenbytheperiodictable.Itisdeterminedbythechargeofthenon‐metalandthesubscriptsthatappearintheformula.

a) iron(II)chloride

b)

c)

d)

e)

f)

g)

h)

I)

j)


Polyatomic Ions

Polyatomic ions consist of a group of atoms combined together that ___________________ with an

overall _______________.

Most polyatomic ions have a __________________, which means they behave as _____________.

This means that they are always written last in the formula.

The one exception: ammonium ion

When writing the formula for compounds containing more than one of a polyatomic ion, the

_________________________________________________________________.

Examples:

Compound Name Formula

barium hydroxide

iron (III) carbonate

copper (I) permanganate

Try these examples yourselves: 1. Ca(OH)2

2. K2CrO4

3. Co(NO2)2

4. Zn(BrO3)

5. ammonium phosphide

6. magnesium borate

7. barium carbonate

8. ammonium nitrate

9. potassium nitrite

10. calcium phosphate

**A useful Chart to help

you with Naming Ionic

Compounds is on page 50

of the notepack!

Assignment: 1. Complete above Examples 2. Do the Practice Problems on pg 46 of

text 3. Do pg 48‐49 of notepack

REMEMBER … Use the same general procedure as

you did with binary ionic compounds but …

• … when you need more than one polyatomic

ion …

• … USE BRACKETS!

3 3 ( )Au NO

s

4 4 ( )3NH PO s

2 2 7 ( )K Cr O s


IonicCompounds‐‐‐PolyatomicIons

Thenamesandchargesofpolyatomicionscanbefoundinlistsandneednotbememorized.Itisagoodidea,however,togettoknowthemorecommononesintroducedinthepracticebelow.Remembertoformthenamebycombinethepositiveandnegativeion:

Name=positiveion+negativeion

COMBINE Ions Formula Name

iron(II)&nitrate iron(II)nitrate

aluminum&nitrate aluminumnitrate

sodium&sulfate

lead(IV)&sulfate

magnesium&carbonate

gold(III)&sulfite

zinc&hydrogencarbonate

ammonium&nitrate

copper(I)&phosphate

silver&hydroxide

aluminum&hydroxide

lead(II)&phosphate

potassium&acetate

manganese(V)&sulfate


NamingIonicCompounds

WritetheEnglishnameofeachofthecompoundsgiven.

ChemicalFormula NameofCompound

1.

2.

3.

4.

5.

6.

7.

8.

9.

10.

11.

12.

13.

14.

15.



Hydrates (extra water )

• Hydrates – ___________________________________________________________________________

___________________________________________________________________________

Examples

CaCl2 5H2O

• _________________________________________

MgSO4 7H2O

• _________________________________________

InternationalChemicalFormula IUPACName1. 2. 3. 4. aluminumsulfide5. zincsulfide6. magnesiumiodide7. 8. 9. 10. tin(II)sulfide11. chromium(III)oxide12. iron(II)sulfide13. 14. 15. 16. ammoniumsulfide17. bariumsulfite18. magnesiumhydroxide19. ∘ 7 20. ∘ 4 21. sodiumsulfatedecahydrate22. 23. bismuth(III)sulfate24. lead(II)acetatetrihydrate25.

Assignment: 1. Do pg 51 of notepack


Molecular Compounds

Remember, for ionic compounds, a formula unit is a ratio of the number of ions in a

________________________________________.

____________________________________________________________________

A molecule is ___________________________________________________ together.

Each molecule is independent of the next, and is not part of a lattice.

____________________________________________________ are formed between two non‐metal

elements.

Properties of Molecular Compounds

Not strong enough to take ____________________

Share ___________ of electrons (covalent bond)

3 main types

o ______________________________

o _________________________________ (not aqueous IE acid)

o ___________________ (2 of the same element bonded together)

Covalent Bonds – Sharing Electrons!

Remember that non‐metals need to gain electrons to have a full ______________________.

When non‐metal atoms combine, the only way this can be achieved is if they _____________ their

outer ________________.

Example:

Since electrons are being _________________, there is a strong force of attraction between the two

atoms. This force is a _______________________.

ionic compound

molecular compound

two chlorine atoms one chlorine

molecule

2 ( )Cl g

a pair of shared electrons

an oxygen atom and two hydrogen atoms

a water molecule

two pairs of shared electrons


Molecular Elements

The vast majority of elements exist in nature as _________ atoms. These are called _______________.

_______________________________________ (exist as pairs of atoms), which you must memorize.

_________________________________________________________

There are two polyatomic elements which also must be memorized:

_______________ _______________________

Naming Molecular Compounds

______________________________________________________________

• ______________________________________________________________________ it is assumed to

be a one (don’t have to put mono) Ex. CO carbon monoxide

Try these Problems:

SO2(g)

CS2(g)

N2O3(g)

CCl4(l)

P4O10(s)

Name Formula

carbon dioxide

dinitrogen monoxide

phosphorus trichloride

2 ( )O F g

2 4 ( )N S g

3 ( )SO g

“I Bring Clay For Our New House.”

“And four Paving stones for eight Steps.”

Assignment: 1. Do Problems to the left 2. Do the Practice Problems on pg 49 of

Text

MEMORIZE THIS TABLE!!!!!


Molecular Compounds that contain Hydrogen

The names, formulas and states for the following molecular compounds containing hydrogen must be

memorized!!!!!

water H2O (l)

hydrogen peroxide H2O2 (l)

ammonia NH3 (g)

glucose C6H12O6(s)

sucrose C12H22O11 (s)

methane CH4 (g)

propane C3H8 (g)

methanol CH3OH (l)

ethanol C2H5OH (l)

hydrogen sulfide H2S (g)

Assignment: 1. Complete pg 56‐57 of Notepack 2. Do C&R pg 50 in Text #1‐12 (a,c,e,g,i)



Binary Molecular Compounds

A. Write the correct name for each compound below. Use prefixes to indicate the number of atoms of each

element in the name of the molecular compound. (Remember: The prefix “mono‐“is not used with the name

of the first element.)

1 atom: mono ‐ 2 atoms: di‐ 3 atoms: tri‐ 4 atoms: tetra‐ 5 atoms: penta‐

6 atoms: hexa‐ 7 atoms: hepta‐ 8 atoms: octa‐ 9 atoms: nona‐ 10 atoms: deca‐

1.

2.

3.

4.

5.

6.

7.

8.

9.

10.

11.

12.

13.

14.

15.

16.


B. Write the correct formula for each compound below. Use subscripts to indicate the number

of atoms of each element in the formula (never reduce).

1. chlorine monoxide

2. sulfur hexachloride

3. dinitrogen monoxide

4. nitrogen trifluoride

5. sulfur tetrachloride

6. xenon trioxide

7. carbon dioxide

8. boron trichloride

9. diphosphorus pentoxide

10. phosphorus trichloride

11. sulfur dioxide

12. bromine pentafluoride

13. disulfur dichloride

14. boron trifluoride

15. tetraarsenic decoxide

16. silicon tetrachloride


Naming Acids

All acids will be considered as being dissolved in water, in aqueous solution and this must always be indicated

by placing the subscript (aq) after the acid formula.

Acids can be named in two ways:

i. IUPAC System

The IUPAC system places the word aqueous in front of the name of the acid, named as if it were an

ionic compound.

ii. Classical System

The classical system uses three different rules, depending on what the acid ends in: – ide, – ate, or – ite.

Rule #1

When the name of the negative ion ends in ‐ide the acid begins with the prefix hydro‐ and the stem of the

negative ion is given the ending –ic, in place of –ide. This is followed by the word acid.

Rule #2

When the negative ion ends in –ate the acid name is the stem of the negative ion given the ending –ic , in

place of –ate , followed by the word acid.

Rule #3

When the negative ion ends in –ite the acid name is the stem of the negative ion given the ending –ous , in the

place of –ite, followed by the word acid.

These rules are found on your periodic table: DO NOT MEMORIZE!!!

Examples:

Acid IUPAC Name Classical Name

HBr(aq) Aqueous hydrogen bromide

H2CrO4(aq)

HNO2(aq)

H2SO4(aq)

HI(aq)

HClO2(aq)

Assignment: 1. Complete pg 59‐60 of Notepack 2. Work on Handout


Completethetablebelow.

Note: *Theclassicalnamedoesnotfollowtheruleasthestemisalteredtosuphu

**Theclassicalnamedoesnotfollowtheruleasthestemisalteredtophospo

AcidFormula

ClassicalAcidName IUPACAcidName

hydrobromicacid aqueoushydrogenbromide

aqueoushydrogenchlorite

* aqueoushydrogensulphate

*

hydrocyanicacid

aqueoushydrogenfluoride

carbonicacid

**

**phosphorousacid


Writetheformulasforthesesubstances:1. Hydrogenfluoride

2. Lead(II)sulphate

3. Magnesiumchloride

4. Tin(II)carbonate

5. Aluminumhydroxide

6. Nickel(III)bromide

7. Calciumnitrate

8. Copper(I)oxide

9. Potassiumcarbonate

10. Chromium(III)oxide

11. Bariumhydroxide

12 Potassiumbicarbonate

13 Copper(II)dichromate

14 Zincphosphate

15 Magnesiumiodide

Writethenamesofthesecompounds1 Na2S

2 PbO2

3 CaCO3

4 SnO2

5 HgNO3

6 H2SO4

7 NH4NO3

8 Co3(PO4)2

9 Al2O3

10 CrBr3


Section 2.3 – Properties and Classification of Ionic and Molecular

Compounds

Properties of Ionic Compounds Crystalline solids (made of ions) at room temperature. ________________________________________________ Will retain their crystal shape when broken. ________________________________________________ When dissolved in water, the solutions will conduct electricity. ________________________________________________

Solubility of Ionic Compounds

Solubility ________________________________________________________________________

If an ionic compound dissolves in water it will form an aqueous (aq) solution.

o NaCl(s) ________________________ o NaCl(aq) ________________________________

If an ionic substance is not soluble in water a solid _____________________ is formed.

All ionic compounds are soluble to some extent, but their solubility is so low that they form precipitates anyway.

Predictions of Solubility To predict whether or not a particular combination of ions will form a soluble compound or not, the solubility table is used. It is found on the back of the periodic table and in the Text on pg 57. How to use the solubility table:

a. Locate the negative ion at the top of the table. b. Below this, find the positive ion either in the row labeled very soluble or slightly soluble. c. The ions found in the slightly soluble will not dissolve in water and will form a precipitate. d. The ions found in the very soluble row will dissolve in water.


Examples:

Properties of Molecular Compounds

1) __________, liquids or __________ (made up of individual molecules) at room temperature. 2) Low melting and boiling points. 3) _______________________________________________________________________ 4) Only some are soluble in water. 5) _________________________________________________________________________

Special Properties of Water

Water is a ______________________ molecule. This means that one end of the molecule has a slight ________________________ and the other end has a slight _____________________________.

Water molecules have a strong attraction for one another.

The force between the positive end of one water molecule and the negative end of a neighbouring molecule is called a _______________________________.

The polarity of water molecules causes water to have several unique properties:

____________________________________.

Large capacity to absorb heat energy without large changes in temperature.

The density of ice (solid state) is less than the density of water (liquid). Water is the only substance on ____________ that has this property. Water’s unique properties have many implications for the existence life on Earth.

Assignment: 1. A2.3 Check and Reflect pg 61 #1,3,5,6,7‐9

2. Notepack pg 63‐64

Compound Soluble Formulasilver sulfate

calcium sulfide

iron(III) sulfide

sodium hydroxide

ammonium sulfate

silver iodide

calcium sulfate

potassium nitrate

ammonium carbonate

lead(II) iodide



L ine Master 11

All Ionic Compounds

I. Write the formula for each compound. **Be sure to include solubility

1) calcium acetate

2) potassium chloride

3) ammonium carbonate

4) sodium nitride

5) titanium(IV) hypochlorite

6) iron(III) sulfide

7) zinc dichromate

8) platinum(IV) oxide

9) aluminium hydroxide

10) mercury(II) nitrate

11) strontium fluoride

12) tin(IV) hydrogenoxalate

13) calcium peroxide

14) gold(I) sulfate

15) lead(IV) thiocyanate

16) nickel(III) sulfide


II.Write the name of each compound.

17) CsI(s)

18) SnCl4(s)

19) Cr(NO3)3(s)

20) (NH4)3PO4(s)

21) Cu2SO4(s)

22) Mg(H2PO4)2(s)

23) Na2S2O3(s)

24) AgClO3(s)

25) Zn(OH)2(s)


Science 10 – Section 2.1 – 2.3 Review

1. Review Bohr diagrams on page 35. Draw a Bohr diagram for:

a) a neon atom b) an oxide ion c) a potassium atom d) a magnesium ion 2. Draw energy level diagrams for: a) phosphorus b) nitrogen c) helium

5. Complete the table below

6. Complete the table below

Name of isotope Mass Number

Atomic Number Number of protons

Number of neutrons

Example: carbon‐12 12 6 6 6

Example: oxygen‐18 18 8 8 10

oxygen‐16

oxygen‐17

64 29

35 45

4 2

Element Number of protons

Number of electrons in an

atom

Number of electrons in an

ion

Charge

Example: oxygen

8 8 10 2‐

Example: sodium

11 11 12 1+

lithium 1+

magnesium

6 3+


Section 2.4 – Acids and Bases Acids An acid is a compound that contains hydrogen and dissolves in water to form a solution that has a pH less than 7. Examples include vinegar, stomach acid etc. Properties of Acids

1) Acids have a _____________ taste. 2) Acids do not have a _________________ feel. 3) Acids react with metals to form hydrogen gas. 4) Acids turn litmus paper __________. 5) Acidic solutions conduct electricity. This is because all acids dissociate

(separate into ions) when they ________________.

Bases A base is a compound that dissolves in water to form a solution that has a pH more than 7. Example: Milk of Magnesia. Properties of Bases

1) Bases have a ______________ taste. 2) Bases have a __________________ feel. 3) Bases do not react with metals. 4) Bases turn litmus paper ___________. 5) Basic solutions conduct electricity. They also dissociate into ions.

*Many basic compounds contain the ________________ ion (OH‐).

The pH Scale The _____________________ is a logarithmic scale.

This means that an decrease in one pH value means a 10 times increase in acidity.

pH means “power of hydrogen.”

Assignment: 1. Pg 67‐68 in Notepack

2. Check and Reflect pg 69

(1,3,4,6,8)



L ine Master 15

Classifying and Naming Compounds

Ionic compounds begin with a metal or the ammonium ion. Molecular compounds contain only non-metals. Acids begin with H or end with COOH.

1. Classify each of the following as an ionic compound, a molecular compound, or an acid. Name each one.

Type Name

a) NaCl(s)

b) N2O(g)

c) HCl(aq)

d) NH4Br(s)

e) KOH(s)

f) CH3COOH(aq)

g) XeF2(s)

h) SCl3(g)

i) NiCl3(g)

j) H3PO4(aq)

k) K2Cr2O7(s)

l) NH4NO3(s)

m) CH3OH(l)

n) Fe2O3(s)

continued...


2. Classify each of the following as an ionic compound, a molecular compound, or an acid. Write the formula for each one.

Type Formula

a) solid sodium sulfate

b) aqueous hydrogen nitrate

c) gaseous sulfur trioxide

d) gaseous dinitrogen trioxide

e) solid manganese(IV) bromide

f) solid ammonium phosphate

g) aqueous hydrogen sulfate

Assignment: 1. Pg 69 in Notepack

2. Read Pgs 70‐75 in Text

3. Do Section 2 Review:

# 1‐22 even, 23‐31 b,d,f,h,j,l, 35,36 b,d,f,h,j,l

4. Prepare for Section 2 Exam


Chemistry Naming and Writing Chemical Formulas

Complete the chart below by writing the appropriate name or chemical formula

NAME CHEMICAL FORMULA

Including state (where applicable)

barium chloride

lead (IV) phosphide

copper (I) sulfide

ammonium phosphate

hexaphosphorus tetrasulfide

barium hypochlorite

heptasulfur nonaoxide

barium iodide

dinitrogen tetrasulfide

uranium (IV) sulfide


Unit A – Section 3.0

Section 3.1 – Chemical Changes Evidence of Chemical Change Recall that a physical change _____________________________________________________________. Examples include: changes in ______________________, ______ and _________. A chemical change is a result of a chemical reaction, where ____________________________________. All chemical reactions have the general form:

The atoms (or ions) of the ___________ are rearranged to form the _______________.

The products have ___________________________ than the reactants and _______________ into or out of the system

Evidence:

1. _______________________

One or more of the products has a different colour than the reactants.

2. ________________________

If one of the products is a gas, bubbles will appear.

3. ________________________

One of the products is only slightly soluble in water.

4. ________________________

This is often noticed as heat or light absorbed or released. o A chemical reaction that releases energy is exothermic. Ex. Fire o A chemical reaction that absorbs energy is endothermic. Ex. Cold Packs

5. _______________________

The products formed during chemical reactions may have a different odour, or no odour at all.

Law of Conservation of Mass Recall that the law of conservation of mass states that the total number of atoms in the

______________ is __________ to the total number of atoms in the ___________.

Bonds between atoms ___________________________________________________________ ______________________________________________________________________________

The atoms of the reactants are simply ______________________ to form the products

Assignment: 1. Read pgs 78 – 85 in Text

2. Chk & Reflect pg 85 # 3, 5,10,11


Section 3.2 – Writing Chemical Equations Chemical equations use _________ or _________________ to show what happens during a chemical reaction. The simplest form of chemical equation is a __________________________

How to write formula equations

1. ____________________________________________________________________________________

2. ____________________________________________________________________________________

3. ____________________________________________________________________________________

4. ____________________________________________________________________________________

Equations must be balanced!! **This means we must always have the same number of each type of atom on both sides of the equation

1) Determine the correct chemical formula for all reactants and products.

_______________________________________________________

_______________________________________________________

_______________________________________________________

2) Balance ___________ 3) Balance ___________ 4) Balance ___________ 5) Balance ___________ 6) Recount all ________ 2) If every coefficient will reduce, rewrite the whole equation using the simplest ratio of coefficients.

Ex. Solid magnesium reacts with aqueous hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas


Balance by inspection **Balance by adding a coefficient to the front of the chemical formula. Coefficients must be _________ numbers. Do not change subscripts in chemical formula. Do not place _______________________ atoms or ions in a formula. Number of polyatomic ions must be __________________________ of the equation

Examples:

KI(aq) + Cl2(g) KCl(aq) + I2(s)

NH3(g) + O2(g) N2(g) + H2O(l)

KClO3(aq) → KCl(aq) + O2(g)

Al(s) + H2SO4(aq) → Al2(SO4)3(aq) + H2(g)

Hg(OH)2(s) + H3PO4(aq) Hg3(PO4)2(s) + H2O(l)

CuO(s) + NH3(g) Cu(s) + H2O(l) + N2(g)

NH3(g) + O2(g) N2O4(g) + H2O(g)


chlorine + magnesium iodide magnesium chloride + iodine

sodium chloride + sulfuric acid hydrochloric acid + sodium sulfate

potassium nitrate decomposes into potassium nitrite and oxygen

bismuth (III) nitrate + calcium iodide bismuth (III) iodide + calcium nitrate

iron (III) oxide reacts with carbon monoxide to produce iron and carbon dioxide

Assignment: 1. Do pg 74‐75of notepack

2. Check and Reflect pg 90 #2,6‐9


Writing Formula Equations from Word Equations In the following exercises, recall that many of the non-metal elements exist as molecules, such as H2(g) or S8(s). Refer to Table A2.9 on page 48 of Science 10 for the chemical formulas of molecular elements.

1. Rewrite the following word equations as formula equations and then balance them:

a) solid sodium metal reacts with chlorine gas to produce solid sodium chloride

b) solid potassium metal reacts with oxygen gas to produce solid potassium oxide

c) hydrogen gas reacts with oxygen gas to produce liquid water

d) solid potassium chlorate decomposes into oxygen gas and solid potassium chloride

e) solid aluminium oxide is decomposed into solid aluminium and oxygen gas

f) mercury(II) sulfide is decomposed into liquid mercury and solid sulfur

g) aqueous cobalt(III) nitrate reacts with solid zinc to produce aqueous zinc nitrate and solid cobalt.

h) fluorine gas reacts with aqueous lead(IV) iodide to produce aqueous lead(IV) fluoride and solid iodine

i) aqueous gold(III) bromide reacts with solid silver metal to produce solid silver bromide and solid gold metal

j) aqueous sodium sulfate reacts with aqueous strontium hydroxide to produce aqueous sodium hydroxide

and solid strontium sulfate


k) aqueous thallium(I) hydroxide reacts with aqueous magnesium bromide to produce solid magnesium

hydroxide and solid thallium bromide

l) methane gas reacts with oxygen gas to produce carbon dioxide gas and water vapour

Balancing Formula Equations 1 Balance the following chemical equations:

a) _____ Na(s) + _____ O2(g) _____ Na2O(s)

b) _____ Al(s) + _____ Cl2(g) _____ AlCl3(s)

c) _____ N2(g) + _____ O2(g) _____ NO2(g)

d) _____ HI(g) _____ H2(g) + _____ I2(s)

e) _____ NH3(g) _____ H2(g) + _____ N2(g)

f) _____ Al2S3(s) _____ Al(s) + _____ S8(s)

g) _____ BN(s) + _____ Cl2(g) _____ BCl3(g) + _____ N2(g)

h) _____ SnF4(aq) + _____ Cr(s) _____ CrF3(aq) + _____ Sn(s)

i) _____ Mg(s) + _____ HCl(aq) _____ MgCl2(aq) + _____ H2(g)

j) _____ (NH4)3PO4(aq) + _____ CaBr2(aq) _____ Ca3(PO4)2(s) + _____ NH4Br(aq)

k) _____ Pb(NO3)4(aq) + _____ K2Cr2O7(aq) _____ Pb(Cr2O7)2(s) + _____ KNO3(aq)

l) _____ AgClO4(aq) + _____ Na3PO4(aq) _____ NaClO4(aq) + _____ Ag3PO4(s)

m) _____ HCl(aq) + _____ Ca(OH)2(s) _____ CaCl2(aq) + _____ H2O(l)

n) _____ CH3COOH(aq) + _____ Ba(OH)2(aq) _____ Ba(CH3COO)2(aq) + _____ H2O(l)

o) _____ C3H8(g) + _____ O2(g) _____ CO2(g) + _____ H2O(g)

p) _____ C6H14(l) + _____ O2(g) _____ CO2(g) + _____ H2O(g)

q) _____ C3H6OS2(s) + _____ O2(g) _____ CO2(g) + _____ H2O(g) + _____ SO2(g)


Balance the following Equations (and add their states)

1) ____ Na3PO4 + ____ KOH ____ NaOH + ____ K3PO4 2) ____ MgF2 + ____ Li2CO3 ____ MgCO3 + ____ LiF 3) ____ P4 + ____ O2 ____ P2O3 4) ____ RbNO3 + ____ BeF2 ____ Be(NO3)2 + ____ RbF 5) ____ AgNO3 + ____ Cu ____ Cu(NO3)2 + ____ Ag 6) ____ CF4 + ____ Br2 ____ CBr4 + ____ F2 7) ____ HCN + ____ CuSO4 ____ H2SO4 + ____ Cu(CN)2 8) ____ GaF3 + ____ Cs ____ CsF + ____ Ga 9) ____ BaS + ____ PtF2 ____ BaF2 + ____ PtS 10) ____ N2 + ____ H2 ____ NH3

11) ____ NaF + ____ Br2 ____ NaBr + ____ F2 12) ____ Pb(OH)2 + ____ HCl ____ H2O + ____ PbCl2 13) ____ AlBr3 + ____ K2SO4 ____ KBr + ____ Al2(SO4)3 14) ____ CH4 + ____ O2 ____ CO2 + ____ H2O 15) ____ Na3PO4 + ____ CaCl2 ____ NaCl + ____ Ca3(PO4)2 16) ____ K + ____ Cl2 ____ KCl 17) ____ Al + ____ HCl ____ H2 + ____ AlCl3 18) ____ N2 + ____ F2 ____ NF3 19) ____ SO2 + ____ Li2Se ____ SSe2 + ____ Li2O 20) ____ NH3 + ____ H2SO4 ____ (NH4)2SO4


Section 3.3 – Types of Chemical Reactions

Synthesis/Formation ________________________________________________ General Formula: A + B → AB Example: 2H2(g) + O2(g) → 2H2O(g) Try these examples: Write the balanced equation for the formation of lithium oxide from its elements. Write a balanced chemical equation for each unfinished formation reaction: calcium + nitrogen → . . . silver + oxygen → . . .

Decomposition ___________________________________________________________________ General Formula: AB → A + B Example: CaCl2(s) → Ca(s) + Cl2(g) Try these examples: Write the balanced equation for the decomposition of solid magnesium sulfide into its elements. Write the balanced equation for the decomposition of solid nickel (II) chloride into its elements.

Hydrocarbon Combustion ____________________________________________________________________________________________________________________________________________________________________________________ General Formula: CxHx + O2 → CO2 + H2O Example: CH4 (g) + 2 O2 (g) → CO2 (g)+ 2 H2O (g)


Single Replacement __________________________________________________________________________ General Formula: A + BC → AC + B Example: Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)

Try these Examples: Write the balanced equations. Chlorine gas is added to a solution of aqueous nickel (III) bromide and the mixture is stirred. This produces aqueous nickel (III) chloride and liquid bromine. Zinc metal is placed into a solution of silver nitrate and allowed to sit. This produces aqueous zinc nitrate and solid silver metal.

Double Replacement ________________________________________________________________________ General Formula: AB + CD → AD +CB Example: H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + 2H2O(l) Try these Examples: Write the balanced Equations. When aqueous copper (I) nitrate and aqueous potassium bromide are mixed, a precipitate of solid copper (I) bromide forms. Another product also forms. When aqueous aluminum chloride and aqueous sodium hydroxide are mixed, a precipitate of solid aluminum hydroxide forms. Another product also forms.

Assignment (all Textbook): 1. Read pgs 91‐97

2. Pg 93 #1‐3

3. Pg 97 #1‐4 Answers in the back of the book


This table is found in your data booklet!! USE IT!!!

Assignment: 1. Read pgs 98‐105 in text

2. Notepack – Pgs 80‐81

3. Textbook ‐ A3.3 Check and Reflect page 106

#’s 1 – 3 (a,c,e,g,i) #’s 4‐8 all



L ine Master 19

Classifying and Balancing Chemical Equations

A. Classify the following reactions as: formation, decomposition, single replacement, double replacement, or combustion.

1. 2 KClO3(s) 2 KCl(s) + 3 O2(g)

2. 3 ZnCl2(aq) + 2 K3PO4(aq) 6 KCl(aq) + Zn3(PO4)2(s)

3. Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)

4. 2 H2(g) + O2(g) 2 H2O(g)

5. Al(s) + 3 NiBr2(aq) 2 AlBr3(aq) + 3 Ni(s)

6. 4 Al(s) + 3 O2(g) 2 Al2O3(s)

7. 2 NaCl(s) 2 Na(s) + Cl2(g)

8. CaCl2(s) + F2(g) CaF2(s) + Cl2(g)

9. AgNO3(aq) + KCl(aq) AgCl(s) + KNO3(aq)

10. 2 C2H6(g) + 7 O2(g) 4 CO2(g) + 6 H2O(g)

11. N2(g) + 3 H2(g) 2 NH3(g)

12. 2 H2O2(aq) 2 H2O(l) + O2(g)

13. (NH4)2SO4(aq) + Ba(NO3)2(aq) BaSO4(s) + 2 NH4NO3(aq)

14. MgI2(aq) + Br2(l) MgBr2(aq) + I2(s)

15. SO3(g) + H2O(l) H2SO4(aq)


B. Balance and classify each chemical equation below.

1. ____ CH4(g) + ____O2(g) ____CO2(g) + ___H2O(g)

2. ____ Na(s) + ____I2(s) ____NaI(s)

3. ____ NaOH(aq) + ____H2SO4(aq) ____Na2SO4(aq) + ____ H2O(l)

4. ____ Fe(s) + ____O2(g) ____Fe2O3(s)

5. ____ Pb(NO3)2(aq) + ____K2CrO4(aq) __ PbCrO4(s) + __KNO3(aq)

6. ____ S8(s) + ____O2(g) ____SO3(g)

7. ____ C3H5(NO3)3(s) __CO2(g) + ____N2(g) + __H2O(l) + ___O2(g)

8. ____ Fe(s) + ____CuCl2(aq) ____FeCl2(aq) + ____Cu(s)

9. ____ C3H8(g) + ____ O2(g) ____CO2(g) + ____H2O(g)

10. ____ CaSO4(aq) + ___AlBr3(aq) ___CaBr2(aq) + ___Al2(SO4)3(s)


Section 3.4 – The Mole

What is the mole? • It is a quantity that measures _________________________________. • Normally we use terms to measure quantities:

Example: 1 pair = 2 items, 1 dozen = 12 items

• The mole ____________________________________________________________________________ • The number is 6.02x1023 and was named after the person who discovered it: Avogadro (hence –

Avogadro’s Number = NA) •

If you had 1 mol of paper how high of a pile would you have? Discovered?

• It is the number of atoms found in 12 grams of carbon‐12 (specific isotope) Carbon‐12 was picked so _________________________________________________________ Carbon 12 is also easily acquired

Why do we use the mole?

• ___________________________________________________________________________________. • We are lucky because for balanced equations it is the ratio of the different species that is important.

Example: 1C6H12O6 + 6O2 6H2O + 6CO2

• The 1:6:6:6 ratio could be 2:12:12:12 or 12:24:24:24 or 1dz:6dz:6dz:6dz or 1mol:6mol:6mol:6mol

How does this help me?

By using a quantity called Molar mass we can quickly count the number of atoms of a given substance.

Molar Mass • ____________________________________________________________________________________ • Unit: g/mol (__________________________________) • Formula symbol “M” • It is found by _____________________________ the total weight of its parts

(elements). • Individual molar mass is found on the ______________________________

(Same number as amu’s but those are for individual atoms) Molar Mass of Compounds

____________________________________________________________________ Calculate the molar masses of the following compounds: a) Water H2O(l)

H = 1.01 g/mol x 2 = 2.02 g/mol O = 16.00 g/mol x 1 = 16.00 g/mol 18.02 g/mol

What this means is that if you counted 6.02x1023 molecules of water it would weigh 18.02g.

1 mole = 6.02 x 1023

Molar masses are

always to two decimal places!


Try these examples:

1. Na2SO4(aq) Na = 22.99 g/mol x 2 = 45.98 g/mol S = 32.07 g/mol x 1 = 32.07 g/mol O = 16.00 g/mol x 4 = 64.00 g/mol

142.05 g/mol

2. (NH4)2SO4

N = 14.01 g/mol x 2 = 28.02 g/mol

H = 1.01 g/mol x 8 = 8.08 g/ mol

S = 32.07 g/mol x 1 = 32.07 g/mol

O = 16.00 g/mol x 4 = 64.00 g/mol

132.17 g/mol

Calculate the molar mass for each of the following (on a separate sheet of paper.

1. NO 2. H2O 3. NH3 4. CO2 5. CH4 6. AgNO3 7. Ca(OH)2 8. Al(NO3)3 9. FeCl3 10. SnC2O4 11. Sn(C2O4)2 12. (NH4)3PO4 13. CH3COOH 14. CH3CH2CH2CH3 15. Ni(H2O)2(NH3)4Cl2 16. Al2(SO4)3 17. Co3(AsO4)2 ● 8H2O 18. Pb(C2H3O2)2 ● 3H2O 19. MgSO4 ● 7H2O 20. KAl(SO4)2 ● 12H2O

Assignment: 1. Complete problems to the left.


Calculatingformoleandmass: Inordertoproperlydeterminesubstancequantitiestouseinachemistrylab,youmustdetermine

thetotalnumberofmolesorthetotalmassrequiredofthatsubstance. Onceyouhavedeterminedthemolarmass,youknowthe#ofgramsin1moleofasubstance(ie.

Cl2is70.90g/mol)

Theformulathatweuseis: Where:

n=numberofmoles(mol) m=mass(g) M=molarmass(g/mol)

**Youwon’thavetomemorizethisformulaasitcanbefoundonpg2ofyourdatabookletManipulatingtheformula:

1. Stepstosolvingconversionproblems:2. Determinewhatyouarelookingfor(massormoles)3. Calculatethemolarmassofthesubstance4. Manipulatetheformulatosolvefortheunknown5. Insertvaluesincorrectplaces6. Solve.

Formula for Mass Formula for moles

PracticeProblems

1. Whatmassisrequiredifyouhave5.0molesofNaCl(s)?

2. Howmanymolesarein25gofCH4?

3. Calculatethemassof2.7molesofwater.


MoleandMassCalculations(Extrapractice)1. Calculatethemassofthefollowing:

a) 1.00moleofNH4Clb) 4.50moleofNH4Clc) 3.25molofPCl3d) 0.00355moleofNa2HPO4e) 0.0125molofXeF4

f) 2.60moleofCH3CH3g) 3.25x102molofNH3h) 7.90x10‐4molofH2SO3i) 1.00x10‐3molofNaOHj) 1.75x10‐4molofFe

2. Calculatethenumberofmolesinthefollowing:

a) 17.0gofH2SO4b) 91.5gofH2Oc) 53.0gofCd) 0.125mgofCuSe) 4.50kgofCH4

f) 225gof(NH4)2SO4g) 55.2mgofCl2h) 128.2gofSO2i) 2955kgofAgj) 0.0845gofKMnO4

3. Calculatethenumberofmolescontainedinthefollowing

a) 7.50x1021moleculesofHNO3b) 425mgofCa(OH)2c) 4.25x1012moleculesofFe2O3d) 0.950kgofNaOHe) 5.50x1025moleculesofCCl4

Assignment: 1. Check & Reflect p 112 (5‐11) in text

2. Section Review P 113 (1,2,5‐8,9a,c,e, 10 a,c,e,

12,13a,c,e,13a,c,e) – Textbook

3. Prepare for Chemistry Unit Exam

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